**2. Crystal engineering**

When the properties of a molecule are tailored in a conventional way by creating or removing covalently bound functional groups or active sites on a molecule, the physical or chemical properties obtained are solely the intrinsic properties of the molecule designed. An alternative approach to modifying the functionality of molecular material is to link molecular units together to create coordination polymers or extended molecular systems. In such systems the interactions between the building units give rise to new properties that do

© 2012 Haukka et al., licensee InTech. This is an open access chapter distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited. © 2012 Haukka et al., licensee InTech. This is a paper distributed under the terms of the Creative Commons Attribution License (http://creativecommons.org/licenses/by/3.0), which permits unrestricted use, distribution, and reproduction in any medium, provided the original work is properly cited.

not exist in the building block molecules. This is the very essence of crystal engineering. Desiraju has stated that "crystal engineering is the rational design of functional molecular solids" and defines crystal engineering as "the understanding of intermolecular interactions in the context of crystal packing and in the utilization of such understanding in the design of new solids".[1] In other words, the goal is to create functional systems by assembling molecular units into extended molecular structures. Over the past few decades vast numbers of papers (Fig. 1) and textbooks have been published on this topic.[2–4]

**Figure 1.** The number of crystal engineering publications since 1985 (ISI WoK, March 2012, topic ="crystal engineering"). These figures include only those publications whose topic includes "crystal engineering". The true number in this field including all related publications is much larger.

As mentioned above, bringing molecules together in a predictable way requires that the intermolecular forces are directional and strong enough to maintain a certain molecular architecture. Non-covalent interactions such as hydrogen bonds, halogen bonds, ππ interactions, metallophilic interactions, and agostic interactions all have directionality to at least some extent. Especially hydrogen bonds, halogen bonds and ππ interactions are relatively strong electrostatic forces with strong directionality. The bond energies of very strong hydrogen bonds range between approximately 65 and 170 kJ/mol, strong bonds between 15 and 65 kJ/mol, and weak hydrogen bonds around 15 kJ/mol or less.[5] The ππ interactions are somewhat weaker with interaction energies up to 50 kJ/mol.[6,7] The strength of the halogen bonds is comparable with the hydrogen bonds ranging between weak (ca. 5 kJ/mol) to strong (180 kJ/mol) contacts.[8] In addition to strength and directionality, the third requirement is that the intermolecular interactions should be selective. If the molecular building blocks contain different types of active sites the contacts must be predictable. Aakeröy et al have shown that even closely related interactions such as hydrogen bonds and halogen bonds can be used side-by-side in the same structure in a hierarcial way to build predictable molecular assemblies.[9,10] The challenge in this kind of combination is that the halogen bond donor (typically iodine or bromine) can interact not only with the halogen bond acceptor (electron-pair donor) but also with the hydrogen bond donor. An example of the coexistence of halogen bond and hydrogen bond is shown Fig. 2. In this example two-point N-HO hydrogen bond contacts and one point IN halogen bond contacts between the 2-aminopyrazin-1-ium and 2,3,5,6-tetrafluoro-4-iodobenzoate are used to build linear chain structure.[10]

**Figure 2.** Halogen bond and hydrogen bond contacts in the linear chain assembly of 2-aminopyrazin-1 ium and 2,3,5,6-tetrafluoro-4-iodobenzoate.[10]

Building predictable assemblies and extended molecular systems is possible only after the very natures of the different types of interactions are understood. When the potential and limitations of these contacts are recognized, they serve as a versatile toolbox for crystal engineering. The following sections will focus on one of the "new" intermolecular contacts i.e. halogen bonds. In fact this is not a new discovery. The first observations about halogen bonds were published as early as 1863. This intermolecular force was, however, almost forgotten for years. But because of the interest in crystal engineering it was "rediscovered" and for the past decade it has become topic of growing interest.

## **3. Halogen bonds (XB)**

144 Recent Advances in Crystallography

not exist in the building block molecules. This is the very essence of crystal engineering. Desiraju has stated that "crystal engineering is the rational design of functional molecular solids" and defines crystal engineering as "the understanding of intermolecular interactions in the context of crystal packing and in the utilization of such understanding in the design of new solids".[1] In other words, the goal is to create functional systems by assembling molecular units into extended molecular structures. Over the past few decades vast

numbers of papers (Fig. 1) and textbooks have been published on this topic.[2–4]

**Figure 1.** The number of crystal engineering publications since 1985 (ISI WoK, March 2012, topic ="crystal engineering"). These figures include only those publications whose topic includes "crystal engineering". The true number in this field including all related publications is much larger.

As mentioned above, bringing molecules together in a predictable way requires that the intermolecular forces are directional and strong enough to maintain a certain molecular architecture. Non-covalent interactions such as hydrogen bonds, halogen bonds, ππ interactions, metallophilic interactions, and agostic interactions all have directionality to at least some extent. Especially hydrogen bonds, halogen bonds and ππ interactions are relatively strong electrostatic forces with strong directionality. The bond energies of very strong hydrogen bonds range between approximately 65 and 170 kJ/mol, strong bonds between 15 and 65 kJ/mol, and weak hydrogen bonds around 15 kJ/mol or less.[5] The ππ interactions are somewhat weaker with interaction energies up to 50 kJ/mol.[6,7] The strength of the halogen bonds is comparable with the hydrogen bonds ranging between weak (ca. 5 kJ/mol) to strong (180 kJ/mol) contacts.[8] In addition to strength and directionality, the third requirement is that the intermolecular interactions should be selective. If the molecular building blocks contain different types of active sites the contacts must be predictable. Aakeröy et al have shown that even closely related interactions such as hydrogen bonds and halogen bonds can be used side-by-side in the same structure in a hierarcial way to build predictable molecular assemblies.[9,10] The challenge in this kind of combination is that the halogen bond donor (typically iodine or bromine) can interact not The definition of halogen bond is not as well established as the definition of hydrogen bond although these interactions have a lot of similarities. Both contacts are electrostatic intermolecular interactions involving an electron donor and an electron acceptor. In hydrogen bonds D-H acts as a hydrogen bond donor, i.e., the electron acceptor. The hydrogen bond acceptors are, then, electron donors such as oxygen or nitrogen atoms (Fig. 3).

**Figure 3.** Comparison of the hydrogen bond (top left) and the halogen bond (bottom left). (D = donor, A = acceptor). Classification of the halogen bonds based on the geometry (right).

Because of the similarities involved, the same terminology has also been adapted in halogen bonds. The halogen in D-X acts as the halogen bond donor (electron acceptor). While electron donors such as nitrogen, oxygen, sulfur etc. act as the halogen bond acceptors (Fig. 3). The key to the halogen bonds is the polarizability of the halogen atom. Therefore, the strongest halogen bonds are formed by the most easily polarizable halogens, and the strength of the halogen bonds typically decreases in the order I > Br > Cl > F.

Halogen bonds are commonly defined as electrostatic interactions between Lewis acids (the halogen atom) and neutral or anionic Lewis bases and abbreviated as XB, where X refers to the halogen and B the Lewis base.[11] The strong directional preferences of a halogen bond arise from the tendency to maximize the main two directional attractive contributions to the interaction energy i.e. electrostatics and charge transfer. These, in turn, minimize the exchange repulsion that is also strongly directional. Optimizing the electrostatic and charge transfer aspects have been successfully used in designing of drugs, liquid crystals, organic semiconductors, magnetic materials, nonlinear optical materials, and templates for solid synthesis.[12–16] Conventionally, halogen bonds have been divided into two classes, TypeI and Type II (Fig. 3), based solely on the bonding geometry.[4] A few theories and concepts have been proposed for rationalizing the XB in greater detail. The most familiar one is the σhole theory. Other theories such as the lump-and-hole theory and the concept of amphoteric halogen bonds have been used to cover the "blind spots" in the σ-hole theory.

#### **4. σ-hole theory**

In most cases, the σ-hole theory has successfully explained the contradictory nature of halogen bonding. Conventionally covalently bonded halogens are seen as negatively charged entities. How, then, is it possible that they can participate in inter-atomic interactions as electron acceptors? In the σ-hole theory the σ-holes are defined as regions of positive electrostatic potential on the outer sides of halogen atoms, centered close to the extension of the halogen atoms' covalent bonds (Fig.4).[17]In general three factors determine the σ-hole's presence or absence and their magnitudes: a) the polarizability of the halogen atom, b) its electronegativity, and c) the electron-withdrawing power of the remainder D of the D-X molecule.[17] When the halogen is more polarizable and has lower electronegativity, the potential of the σ-hole can become more strongly positive. So the positivity of the σ-hole increases in the order F ‹ Cl ‹ Br ‹ I.

**Figure 4.** Regions of concentrated negative electrostatic potential (blue) and regions of depleted potential (red) on pentafluoroiodobenzene.

Apparently, the σ-hole determines the existence and, the strength of the halogen bonding. Since the σ-hole is located on the extension of the covalent bond along the D-X axis, it generates the directional preferences of the halogen bonding. When the halogen atom in D-XA acts as the halogen bond donor, the D-XA angle is close to 180° (Fig. 3 and Fig. 4). If the halogen acts as the halogen bond acceptor (electron donor), the angle is close to 90° because the electron density around the halogen atom is concentrated at an angle of 90° from the D-X bond (Fig. 3).
